Chapter 4

Chapter 4: Chemical Bonding and Molecular Geometry

4.1 Ionic Bonding

• Formation of Ions

  • Cations: Positive ions formed when neutral atoms lose electrons from the valence shell.

  • Anions: Negative ions formed when neutral atoms gain electrons.

  • Ionic Compounds: Composed of cations and anions held together by ionic bonds (electrostatic attraction).

• Properties of Ionic Compounds:

  • Crystalline structure, rigid, brittle, high melting/boiling points.

  • Poor electrical conductors in solid state; good conductors when dissolved in water or melted.

  • Example: Sodium chloride (NaCl) formed from sodium cation (Na+) and chloride anion (Cl–).

• Formation of Binary Ionic Compounds:

  • Consist of a metal (cation) and a nonmetal (anion).

  • Metals lose electrons easily, nonmetals gain electrons to fill valence shells.

  • Example: Formation of NaCl from sodium and chlorine.

• Charge Balance in Ionic Compounds:

  • Total positive charges from cations must balance total negative charges from anions.

  • Example: Aluminum oxide (Al2O3) contains two Al3+ and three O2– ions.

4.2 Covalent Bonding

• Covalent Bonds:

  • Formed by mutual attraction of two atoms for shared electrons.

  • Occur between nonmetals with similar electronegativities.

• Characteristics:

  • Generally lower melting/boiling points compared to ionic compounds.

  • Many are gases or liquids at room temperature; softer than ionic solids.

  • Poor electrical conductivity in any state.

• Electronegativity:

  • Measure of an atom's tendency to attract shared electrons.

  • Electronegativity difference determines bond polarity.

• Types of Covalent Bonds:

  • Pure Covalent Bond: Electrons shared equally (e.g., H2, Cl2).

  • Polar Covalent Bond: Electrons shared unequally, leading to dipole formation (e.g., HCl).

4.3 Chemical Nomenclature

• Naming Ionic Compounds:

  • Include the cation name followed by anion name with the suffix -ide (e.g., NaCl = sodium chloride).

  • Metals that can form multiple charges have their charge indicated by Roman numerals (e.g., Iron(II) chloride for FeCl2).

• Naming Molecular Compounds:

  • Use prefixes to denote the number of atoms in the molecule (e.g., CO2 = carbon dioxide).

• Naming Acids:

  • Binary acids derived from hydrogen and one nonmetal: change nonmetal's prefix to hydro-, add -ic, followed by "acid" (e.g., HCl = hydrochloric acid).

  • Oxyacids named based on the polyatomic anion present; e.g., from sulfate (SO4^2−) to sulfuric acid (H2SO4).

4.4 Lewis Symbols and Structures

• Lewis Symbols:

  • Show valence electrons as dots around the elemental symbol.

• Lewis Structures:

  • Depict bonding between atoms in a molecule or ion.

  • Follow the steps to draw structures for easy visualization of atoms, bonds, and lone pairs.

4.5 Formal Charges and Resonance

• Formal Charge:

  • Hypothetical charge calculated based on electron assignment in Lewis structures (formula: # valence electrons - # lone pair electrons - 1/2 # bonding electrons).

• Resonance:

  • More than one valid Lewis structure for a molecule exists due to different electron arrangements.

  • Actual structure is a resonance hybrid of all forms present.

4.6 Molecular Structure and Polarity

• Predicting Geometry:

  • Use VSEPR Theory to predict molecular shapes based on electron pair arrangements.

• Molecular Polarity:

  • Determined by bond polarities and molecular geometry; if dipoles cancel out, the molecule is nonpolar.

  • Polar molecules can align in electric fields and dissolve in polar solvents.